CH221 CLASS 1
CHAPTER 1: STRUCTURE AND BONDING
Synopsis. The first class of Organic Chemistry 1 begins by tracing a brief history
of the subject, followed by an overview of the status of modern organic chemistry.
Next, the bonding in organic molecules is discussed, starting with early valence
ideas and ending with the more recent valence bond ideas. The next class deals
with an atomic and molecular orbital description of bonding.
A Brief History of Organic Chemistry
The science of chemistry evolved slowly from the art of alchemy and emerged as
a separate science during the late 1700s. Very early in the development of
chemistry, it was recognized that there were (at that time) mysterious differences
between substances found in plants and animals and those obtained from the
earth as minerals. The former, generally more difficult to isolate and handle, and
generally less stable, were called organic and the latter were known as
inorganic.
The Vital Force
Most chemists in the early 1800s believed that the differences between the two
classes of chemical substances described above was that organic compounds
contained a “vital force”, because they originate in living organisms. Furthermore,
it was widely believed that this vital force prevented organic compounds from
being synthesized or transformed by ordinary laboratory procedures. It gradually
became evident that organic substances can be transformed in the laboratory in
much the same way as inorganic substances. Two major pieces of evidence for
this came from the work of Michel Chevreul, in France and Friedrich Wöhler, in
Germany, as summarized on the next page. Of course, unknowingly, mankind
had manipulated the transformation of organic substances for centuries, as in
brewing and winemaking, baking, dyeing and tanning, to mention just a few
activities.
aqueous
alkali
Animal fat
aqueous
acid
Soap
soap +
glycerin
(glycerol)
fatty acids
[Chevreul, 1816]
heat
NH4+ OCNammonium cyanate
(NH2)2CO
urea
[Wohler, 1828]
Below is a translation from the German of Wöhler’s famous 1828 publication.
ON THE ARTIFICIAL PRODUCTION OF UREA
by F. Wöhler
Annalen der Physik und Chemie, 88, Leipzig, 1828
In a brief earlier communication, printed in Volume III of these Annals, I stated that by the
action of cyanogen on liquid ammonia, besides several other products, there are formed oxalic
acid and a crystallizable white substance, which is certainly not ammonium cyanate, but which
one always obtains when one attempts to make ammonium cyanate by combining cyanic acid
with ammonia, e.g., by so-called double decomposition. The fact that in the union of these
substances they appear to change their nature, and give rise to a new body, drew my attention
anew to this subject, and research gave the unexpected result that by the combination of
cyanic acid with ammonia, urea is formed, a fact that is noteworthy since it furnishes an
example of the artificial production of an organic, indeed a so-called animal substance, from
inorganic materials.
By 1850, the vital force theory had been dropped and the subject of organic
chemistry began to develop very rapidly, especially in the laboratory sense: the
number of new organic substances prepared started on a steep upward path.
From the point of view of bonding and structure, development was a little
slower: this is discussed later.
The Subject of Organic Chemistry
It was soon discovered that all organic substances contain the element carbon
and it is now evident that all organic molecules possess at least one carbon
atom and, indeed, the huge majority possess more than one carbon atom. So,
organic chemistry can be described as the study of the chemistry of
carbon compounds. This definition holds irrespective of whether the
compound is natural or not, although a great deal of the modern emphasis of
organic chemistry is on natural products and the relationships between organic
chemistry and life sciences and medical sciences are strong ones.
Nowadays (2005), there are nearly 20 million organic compounds in the CAS
registry of the American Chemical Society and thousands of new ones are
synthesized each week in laboratories all over the world. In fact, the
compounds of carbon easily outnumber the compounds of all other elements.
The reasons why carbon is able to form such a large number of compounds,
with a vast range of size and complexity, are to be found in the position of
carbon in the periodic table. Carbon is a second row (period) element: it is the
first element of group 4A. As such, it is endowed with the following properties:
1. Carbon is able to form 4 strong covalent bonds. D-orbitals
cannot be used to expand the valence beyond 4, as in the case
of some silicon compounds. Thus alkanes (CnH2n+2) are much
more stable than silanes (SinH2n+2) and carbon tetrachloride
(CCl4) is much less reactive than its silicon analog (SiCl4).
2. Carbon is able to form long chains and rings. This is a
feature of other group 4 elements, but this ability diminishes
rapidly as one proceeds down the group.
3. Carbon, like other elements of period 2 (especially N and O),
is able to form multiple bonds.
Modern organic chemists are extremely skilled at synthesizing new organic
chemicals in the laboratory. Medicines, polymers, dyestuffs, food additives,
agrochemicals, including pesticides, and a huge range of other substances are
prepared (at least initially) in the laboratory. Organic chemistry touches the
lives of everyone – from the good (for example, medicines and polymers)
through the bad (for example, pollution) to the ugly (for example, nerve gases).
Organic chemistry as a subject interacts with many other subject areas, within
and without the general subject of chemistry: examples include physical and
computational chemistry (e.g. structure, reactivity, molecular design), inorganic
chemistry (organometallic compounds), biochemistry and medicine (e.g.
metabolic pathways, drug design, drug metabolism), pharmacology,
agricultural science (e.g. pesticides), food science and geology (hydrocarbons).
Some of these relationships are admirably summarized by Homer Simpson:
Atomic Structure and Chemical Bonding
Organic chemistry uses the familiar nuclear model of the atom, with most of
its mass residing in the tiny (10-14 – 10-15 m in diameter) nuclear subatomic
particles, protons and neutrons, at the center and surrounded by electrons at
a distance of about 10-10 m. Hence, the diameter of a typical atom is about 2 x
10-10 m (= 200 pm, 0.2 nm or 2 Å). An atom is described by its atomic
number (Z), which indicates the number of protons in the nucleus and by its
mass number (A), which indicates the total of protons plus neutrons. All
atoms of the same element have the same atomic number, for example, 6 for
carbon, 1 for hydrogen, 8 for oxygen, etc., but they can have different mass
numbers, depending on the number of neutrons in the nucleus. These are
called isotopes and examples include 12C (6 neutrons), 13C (7 neutrons) and
14C (8 neutrons); 1H, 2H (deuterium) and 3H (tritium); 16O, 17O and 18O.
Organic chemistry also uses the wave mechanical model of the atom, which
treats electrons primarily as standing waves. The properties of a particular
electron in an atom are described by a wave equation (e.g. the Schrödinger
equation) and result from a particular solution to this equation, called a wave
function, ψ. A three-dimensional plot involving ψ2 describes the volume of
space around the nucleus where a particular electron is most likely (say 90 –
95%) to be found. This three-dimensional envelope is called an orbital. There
are several types of orbitals, beginning with the simplest s-type to more and
more complex p-, d-, f-, g-,… types. In organic chemistry, because carbon
belongs to period 2, s and p orbitals are the most important, but d orbitals can
be important when carbon is bonded to period 3 elements, such as
phosphorus and sulfur, and beyond. S, p and d orbitals are illustrated below.
An s orbital, found in all periods or shells
A p orbital, found in period 2 (2nd shell)
onwards
A d orbital, found in period 3 (3rd shell)
onwards
Orbitals are arranged in layers or shells of increasingly larger size and higher
energy, as summarized below.
The three different p orbitals in a particular shell are oriented in space along
mutually perpendicular directions (they are orthogonal) and are denoted by p x,
py and pz. Each p orbital has a plane of zero electron density through its
center – this is called an angular node.
Atomic Structure and Electronic Configuration
Orbitals are filled with electrons in order of increasing energy and taking into
account the Pauli exclusion principle and Hund’s rule. The electronic
configurations that result are called ground state configurations. This is
illustrated for carbon, below – see textbook, p 5, for further examples.
It will be seen later that this ground state electronic configuration of carbon
(and other elements) is not normally the one actually used in chemical
bonding!
Development of Chemical Bonding Theory
Although John Dalton (1806) depicted molecules as being made of atoms that
were somehow linked together, information on atomic weights and molecular
composition was limited at that time and ideas about valence were
undeveloped. Thus Dalton represented methane (“carbureted hydrogen”) as
CH2. It was not until the mid 1800s that carbon was described as being
tetravalent in its organic compounds (Archibold Couper and August Kekulé).
Slightly later came ideas on multiple bonding and rings, especially the
aromatic ring of benzene (Kekulé, 1865).
The three dimensional concept of molecules (and hence of covalent chemical
bonding) was developed largely by Joseph Le Bel and Jacobus van’t Hoff in
1874, the latter proposing the tetrahedral nature of tetravalent carbon:
The Nature of Chemical Bonding
Atoms bond together because the resulting molecule is more stable (of lower
energy) than the separate atoms, or smaller molecular units. For example:
The latter energy change is known as the bond energy: it refers to the
energy needed to break one mole’s worth of bonds of that type in that
particular molecule. From an electronic point of view, a filled outermost shell
(valence shell), like that of a noble gas (group 8A) has a special stability.
Elements take part in chemical bonding in order to acquire this filled valence
shell, which contains 8 electrons (but only 2 for hydrogen) and hence is
known as a complete octet. Elements can do this by complete transfer of
electrons (ionic bonding) or by sharing electrons (covalent bonding), the latter
of which is of prime importance in organic chemistry. Gilbert Lewis (1916) was
the first person to describe covalent bonds in terms of shared electrons: he
depicted molecules using electron-dot structures, now called Lewis
structures:
Nowadays, although the ideas of Lewis are fully accepted, line-bond
structures are more conveniently used. Here, the bonding electrons are
shown as lines (one line for each pair) and non-bonding (lone pair) electrons
and unpaired electrons are shown as dots:
Valence Bond Theory
There are two models to describe how electrons are shared in covalent
bonds: the first is known as the valence bond theory and the second,
considered in the next class, is called the molecular orbital theory. The two
models, each with its own strengths and weaknesses, are complementary
and tend to be used interchangeably by organic chemists, depending on the
situations. The valence bond approach is rather more easy to visualize and
hence will be used more often than the molecular orbital method.
According to valence bond theory, a covalent bond is formed when two atoms
approach each other sufficiently closely so that a singly occupied orbital on
one atom overlaps a singly occupied orbital on the other atom. The two
electrons are now paired in the overlapping orbitals and are attracted to the
nuclei of both atoms: they also shield each nucleus from the other. This is
illustrated below for the formation of a hydrogen molecule.
The above process can be described by an energy diagram, which shows
clearly that the hydrogen molecule is more stable than the two hydrogen
atoms by 436 kJ/mol. Conversely, as previously explained, it requires 436
kJ/mol of energy to break the bond into its component atoms and hence this
energy is known as the bond energy:
How closely the nuclei approach each other when atoms form covalent bonds
is governed by a balance between attractive forces and repulsive forces, as
shown on the energy/internuclear distance diagram for the formation of H 2,
overleaf. The energy minimum corresponds to the formation of the covalent
bond, which in turn corresponds to an optimum distance, the bond length. In
particular, if the nuclei approach each other more closely than this distance,
the energy increases rapidly. Thus every covalent bond has both a
characteristic bond energy and bond length. Some examples of
characteristic bond energies and bond lengths are given in the table overleaf.
.
Bond
Average
bond
energy / kJ/mol
Average
length/ pm
C-H
C-C
412
348
109
154
C=C
612
837
360
743
463
484
238
134 (139 in benzene)
120
143
122
96
138
214
CC
C-O
C=O
O-H
C-F
C-I
bond